The Synthesis of a Complex Iron Salt
Introduction:
In this experiment a complex compound containing the elements potassium, iron, carbon, hydrogen, and
oxygen will be synthesized. Carbon and oxygen will be present in the compound as the oxalate ion (C 2O4-2)
whereas hydrogen and oxygen will be present as water. The final product, consisting of emerald green crystals,
may be given the empirical formula KxFey(C2O4)z . nH2O, where the nH2O is called the water of hydration.
This will be the first of a series of five experiments in which the complex salt will be synthesized, and then its
simplest formula (ie. x, y, z, and n) will be determined, using a variety of analytical techniques.
One important factor in any chemical synthesis is the actual quantity of desired product obtained compared to
the theoretical amount predicted on the basis of the stoichiometry of the reaction. The ratio of the mass of the
product obtained to the theoretical quantity, expressed as a percentage, is referred to as the 'percent yield' or
more simply the 'yield'.
There are many reasons why the actual yields are not 100%. Possibly the reaction reaches equilibrium before
going to completion. Maybe the reactants are involved in reactions other than the one that produces the desired
product. Probably some product is lost in crystallizing and separating crystals from the supernatant liquid, etc.
In this experiment, a solution containing 3.20 grams of FeCl3 . 6H2O (MW= 270 grams/mole) will be reacted
with an aqueous solution containing excess K C O .H O to produce K Fe (C O ) . nH O.
2 2 4
Questions:
(1)
x
y
2 4 z
2
Assuming that all of the iron (Fe) originally in FeCl3 . 6H2O ends up in the product,
K Fe (C O ) . nH O, how many moles of product should be obtained?
x
(2)
2
y
2 4 z
2
What additional information would be needed to calculate a 'percent yield'?
To prepare several grams of pure emerald green crystals of KxFey(C2O4)z . nH2O.
Objective:
Apparatus: two 50-ml beakers, one 800-ml or 1000-ml beaker, two 250-ml beakers, beaker tongs, electronic
balance (0.0001-gram accuracy), 9 mm watch glass, vacuum filtration apparatus with Büchner
funnel, medium flow filter paper, small brown bottle with cap (or black film canister)
Chemicals:
acetone
FeCl3.6H2O solution (0.400 gram of FeCl3.6H2O per ml of solution), K2C2O4.H2O, ice,
Safety, Environmental, and Economic Concerns:

Waste solutions may safely be discarded down the drain. Flush with excess water.

Avoid overheating since overheating may result in violent boiling when heating liquids.

ACETONE IS A FLAMMABLE SOLVENT! Extinguish all flames and turn off hot plates in your
work area before using acetone. Makes sure that the acetone bottle is covered when not in use. If a
spillage occurs, make sure all burners and hot plates in the area are turned off, then clean up the
spilled liquid and use large quantities of water to rinse it down the drain.
-2Notes on Experimental Procedure:
1. When a desired product is formed by crystallization from a reaction mixture containing excess reactants
and other products, the crystals are likely to be relatively impure. The crystals can be separated from the
impure solution (often called the "mother liquor") by filtration or decantation. The mother liquor
clinging to the crystals can be removed by washing with an appropriate solvent. However, washing will
not remove impurities occluded within the crystals.
A standard method for purifying a crystalline product is recrystallization. For crystals that are more
soluble in hot solvent than in cold solvent, the recrystallization can be done by dissolving the crystals in
a minimum quantity of hot solvent and then cooling it in an ice bath. The purified crystals can then be
"harvested" by filtration. A second "crop" of crystals can be obtained from the filtrate by evaporating a
fraction of the solvent by heating, followed by cooling the remaining solution in an ice bath. The second
crop of recrystallized product is generally less pure than the first.
2. The wet crystals dry very slowly. The purpose of the acetone is to wash the water off the crystals. The
acetone, which has a high vapor pressure, evaporates quickly, thus leaving the crystals dry. Since
acetone is flammable, make sure there are no flames on the workbench when doing your acetone washes.
3. The product will slowly decompose when exposed to light. Hence the crystals should be stored in a light
resistant sample bottle (brown amber bottle or black film canister).
Experimental Procedure:
Day 1

Obtain (in a clean, dry 50-ml beaker) 8.00 ml of the stock solution of iron. (The stock solution contains
0.667 grams of ferric chloride hexahydrate per ml of solution).

Weigh 12 grams of potassium oxalate monohydrate (K2C2O4.H2O) into a clean, dry 50-ml beaker, using
the Analytical balance (record to 0.0001 gram). Add 20 ml of distilled water to dissolve the salt
(K2C2O4.H2O). Heat on the hot plate & stir to dissolve the salt (take care not to heat too strongly).


Using beaker tongs to handle the hot beaker, pour the hot solution into the beaker containing the iron
(III) chloride solution, and stir.
Cool the solution for 20 - 35 minutes by placing the beaker in a large beaker containing ice and water
(taking care that it does not sink into the ice water). Crystals should form during this time.

After giving the crystals ample time to form, carefully pour off and discard the solvent without removing
any crystals - a process called decantation.

Add 20 ml distilled water to the crystals. Heat gently with stirring to completely dissolve the crystals. If
some dark residue remains undissolved, carefully decant the clear solution into another beaker and
discard the residue.

Cover the beaker with a watch glass and set it in your drawer until the next lab period in order to allow
the crystals to form. If the crystals are allowed to form slowly without being disturbed, large crystals
will be obtained. If the solution is moved, stirred, or disturbed while the crystals are forming, smaller
crystals will result.

Clean a small sample bottle. Allow it to drain & thoroughly dry it with a paper towel.
-3-
Day 2:

Set up a vacuum filtration system and filter the crystals using a Büchner funnel and a clean filter
flask. Be sure to attach a water trap between the filter flask and the aspirator. Make sure the filter
paper is properly "seated".

Wash the crystals twice with ice water (distilled). [Put a distilled water bottle into the ice bucket.]
Use less than 5 ml of ice water for each wash and work quickly to avoid dissolving the product in
the wash water. Finally rinse the crystals twice with 5 ml portions (aliquots) of acetone.

Spread the crystals in the bottom of a clean, dry 250-ml beaker and set it aside to air dry in your
locker.
Day 3:

Using the Analytical balance, weigh to the nearest 0.0001 gram the clean, dry brown bottle.

When the crystals are dry, place them in the pre-weighed brown bottle and weigh again to the nearest
0.1 milligram. Store these crystals in your drawer in the capped amber bottle for use in the future
experiments. A minimum of 3.5 grams of product will be needed in subsequent experiments. If
your yield is less than 3.5 grams, consult with me. Do NOT put the bottle of green crystals in a
desiccator or in the drying ovens for further drying! [Question: Do you know why this would be a
bad technique?]
Raw Data:
1. volume of iron(III) chloride solution used ..................................... ___________________
2. mass of solid K2C2O4.H2O and beaker ....................................... ___________________
3. mass of beaker ............................................................................... ___________________
4. mass of amber bottle & complex crystals ...................................... ___________________
5. mass of amber bottle ...................................................................... ___________________
Calculated Data: (Show your calculations)
1. mass of solid K2C2O4.H2O ................................................... ___________________
2. mass of complex crystals ........................................................ ___________________
3. moles of ferric ion in reactant sample ........................................... ___________________
-4-
Determination of the % H2O in an Iron Oxalato Complex Salt
Introduction:
In the previous experiment a green crystalline product having the formula KxFey(C2O4)z . nH2O was prepared.
In the next few experiments, the percentage composition of the compound will be determined. From that
information, the empirical formula will be derived. Finally the percent yield will be determined.
The green iron oxalato complex is one of a number of solid chemicals that are classified as "hydrates". A
hydrate contains water chemically bound in the solid state so that it is present in the compound in stoichiometric
amounts. Familiar examples of hydrates are Plaster of Paris (CaSO4.1/2H2O), gypsum (CaSO4.2H2O), and
alum [KAl(SO ) .12H O]. The water of hydration of many hydrates can be removed as a gas by heating the
4 2
2
hydrate to a temperature above 100oC for a period of time. The following reaction involving barium chloride
dihydrate (BaCl2.2 H20) occurs when the solid is heated above 100oC.
BaCl2.2H2O(s) --------------> BaCl2(s) + 2 H2O(g)
The percentage of water of hydration in KxFey(C2O4)z . nH2O will be determined in this experiment by heating
a weighed sample of green hydrate in an open container in a drying oven until all of the water of hydration has
been driven off.
KxFey(C2O4)z . nH2O --------------> KxFey(C2O4)z + nH2O(g)
The loss in weight will be set equal to the mass of the water of hydration.
.............................................................................................................................
Objective:
To determine the percentage of water of hydration in the compound KxFey(C2O4)z . nH2O.
Apparatus:
Analytical and electronic top-loader balances, two 50-ml beakers, desiccator, beaker tongs,
crucible tongs, 800-ml beaker, 150 mm watch glass, crucible & cover, drying oven.
Chemicals:
Student-prepared KxFey(C2O4)z . nH2O
Safety, Environmental, and Economic Concerns:
Waste chemicals from this experiment may be safely discarded in the solid waste receptacle in the lab.
...........................................................................................................................
Notes on Experimental Procedures:
1. The mass of the green crystals must be accurately measured both before heating as well as after heating
(the mass loss upon heating). The mass of KxFey(C2O4)z . nH2O can be calculated by subtracting the
mass of the empty crucible and cover from the crucible, cover, and crystals. The mass of the water of
hydration can be calculated by subtracting the mass of the crucible, cover, and crystals after heating
from the mass of the crucible, cover, and crystals before heating. [Continue to dry, cool, and weigh the
sample until 2 consecutive masses agree to within 0.0010 gram of each other.]
-5Notes on Experimental Procedures (cont.):
2. When crucibles & covers are stored in a desiccator or placed in the oven to dry, each crucible & cover
should be placed in a small beaker to guard against tipping over or becoming contaminated on the
outside. Always have crucible lids at least slightly open while they are being heated. The crucible
should be completely covered for long-time storage at room temperature. After they have been dried in
the oven, the crucibles & covers should not be touched by fingers until the experiment is completed due
to the moisture and oils your hands may impart. Handle the crucible & cover with crucible tongs.
3. The most frequent cause of erratic balance readings is failure to have objects at room temperature when
they are being weighed. Convection currents in a closed balance compartment can have a surprising
effect. Be sure that the glass windows on the balance are closed on both sides, by the way.
4. A small inaccuracy in the calibration of a balance is cancelled out in the final results when the mass of
the empty container and the mass of the container plus the substance of interest are measured using the
same balance. Therefore, the same balance should be used for all successive weighings.
Experimental Procedures:
Day 4:

Before beginning this experiment, the green crystal should have been dried at room temperature in your
locked drawer for several days, weighed, and stored in a capped amber bottle.

Mark two crucibles with a marking pen to distinguish one from the other, then place them (along with a
cover) in separate 50-ml (or 100-ml) beakers. Place each container in the drying oven at 110º - 120º C to
dry for at least 25 minutes.

Cool the beakers and crucibles (with covers) in a desiccator for at least 10 minutes.

Weigh both of the crucibles & covers on the analytical balance to the nearest 0.0001 gram.
Days 5, 6, & 7:

Pre-weigh about 1.0 gram of the green crystal into each of the two crucibles (use the top loader). Then
weigh the crucibles, covers, and crystal samples to the nearest 0.0001 gram. Be sure to use the same
analytical balance for all weighings (same balance that was used to weigh the empty crucible & cover).

Place the crucibles, covers, and crystal samples in the 50-ml beaker and heat in the drying oven for two
hours at 110º - 120º C. Do not leave them in for a longer time or heat at a higher temperature because
the crystals may begin to decompose and turn brown. [It will probably be necessary to heat them for
about one hour on the first day, then an additional hour the next day.]

On the third day, return the crucibles, cover, and contents to the oven for an additional 30 minutes of
heating followed by cooling, and then weigh to ascertain whether more water has been driven off.
Repeat this procedure until the crucibles, covers, and remaining solid have a constant mass.

Calculate the loss in mass upon heating for each sample and attribute the mass loss to the water of
hydration. Determine the percent water of hydration for both samples, and compute the average.

After the final weighings, place the amber bottles with the remaining crystals (capped) back into the lab
drawer. (Do not mix the anhydrous crystals with the original crystals!)
-6-
Data Table for Water of Hydration in a Complex Iron Salt
Trial #1
Trial #2
1. mass of crucible, cover, and hydrated crystal sample (in grams)............ __________
__________
2. mass of crucible, cover, & crystal sample after 1st heating (in grams)... __________
__________
3. mass of crucible, cover, & crystal sample after 2nd heating (in grams).. __________
__________
4. mass of crucible, cover, & crystal sample after final heating (in grams). __________
__________
5. mass of empty crucible & cover (in grams)............................................ __________
__________
6. mass of hydrated crystal sample (in grams).......................................... __________
__________
7. mass of anhydrous crystal sample after final heating (in grams)........... __________
__________
8. percentage of water of hydration in the crystal sample........................ __________
__________
9. average percentage of water of hydration .............................................
Show any calculations:
_______________
-7-
Determination of the Empirical Formula of a Complex Iron Salt
Percentage of Potassium & Iron by Ion Exchange
Introduction:
In a previous experiment, the percentage of water of hydration in the green iron oxalato complex salt,
KxFey(C2O4)z . nH2O, was determined. This experiment involves determining both the percentage of
potassium (K) and of iron (Fe). A titration will be performed with a sample of a solution that has passed
through an ion exchange column. This solution will contain a known mass of the iron complex salt.
Ion Exchange: Certain materials called ion exchange resins consist of rather large molecules, which contain
ionizable groups. The resins are solids - insoluble in water and usually granular in nature - which, when added
to water, swell to form a slurry. The ionizable group on the resin ionizes (exchanges ions) in the presence of
water. This process is shown by the following equation for a resin containing a sulfonic acid group (-SO3-H):
R-SO -H + H O -------------> R-SO - H O+
3
2
3
3
where 'R' represents the large insoluble resin molecule to which the sulfonic acid group is chemically bonded,
and H3O+ represents the hydronium ion bound to the resin sulfonate ion. This particular type of resin is called a
cation exchange resin, and the chemical form of the resin shown in the reaction above is called an acid form
resin.
A slurry of resin in water is poured into a vertical glass or plastic column equipped with a porous plug at the
bottom to trap the resin. Excess water is allowed to flow out, and the column becomes filled with the watersoaked resin (see figure below):
resin
porous plug
If an aqueous solution of a salt such as KCl (K+ and Cl- ions) is poured into the resin-filled column, the KCl
solution will displace the solution surrounding the resin and a volume of liquid equal to the volume of KCl
solution added will elute (be washed out) from the bottom of the column. In the process, as the KCl solution
passes down the column, potassium (K+) ions displace (exchange with) hydronium ions, and aqueous HCl
(H3O+ + Cl-) elutes from the column as shown by the following reaction:
R-SO - H O+ + K+ + Cl- ------> R-SO - K+ + H O+ + Cl3
3
3
3
Thus the solution coming out of the column (the elutant) will contain a quantity of hydronium (H3O+) ions
equal to the number of potassium (K+) ions that were added to the column. If enough potassium ions are added,
all of the acid form of the resin will be converted to the potassium form and at that point the resin will become
incapable of exchanging any more hydronium ions for potassium ions. However, it is possible to restore the
resin completely to its acid form by pouring an aqueous solution of HCl into the K+ saturated column. As the
HCl solution passes down the column, the hydronium ion displaces the bound potassium ions, and the reaction
represented by the equation above is reversed, with aqueous KCl eluting from the column.
-8Other cations such as Na+, Li+, Ca+2, etc. will exchange with the resin-bound H3O+ ions in a manner similar to
that of K+ - hence the term cation exchange resin. Anion exchange resins are also available, but will not be
used in this experiment.
Ion exchange resins are widely used in industry & in research laboratories to selectively remove certain ions
from solution. Home water-softening units are packed with a sodium (Na) form of cation exchange resin, which
removes cations such as Ca+2, Mg+2, and Fe+2 that cause water "hardness". These resins can be regenerated
after they become saturated with the above ions by passing an aqueous solution of NaCl through the unit.
------------------------------------------------------------------------------------------------------------------------------
Determining the Percentage of Potassium in KxFey(C2O4)z . nH2O
Using Cation Exchange
When a weighed quantity of KxFey(C2O4)z . nH2O is dissolved in water, the salt dissociates into ions
according to the following equation:
KxFey(C2O4)z . nH2O -------------> x K+
+ Fey (C2O4)z(aq)-x
+ n H3O+(aq)
If this solution is passed down a column containing a cation exchange resin in the acid form, the K+ ions
will replace the resin bound H3O+ ions according to the following equation:
x R-SO3- H3O+ + x K+ + Fey (C2O4)z(aq)-x -----> R-SO3- K+ + x H3O+ + Fey (C2O4)z(aq)-x
Thus for each mole of potassium ion (K+) added to the column, one mole of hydronium ion ( H3O+ ) elutes
from the column.
moles K+ added = moles H3O+ eluted
If the eluted solution is titrated with a standardized (3 sig. fig.) 0.1 M NaOH solution, the moles of H 3O+ in
the solution eluted from the column (hence the moles of K+ added to the column) can be determined.
moles K+ added = moles H3O+ eluted = moles NaOH used in titration
moles NaOH used in titration = [Volume (in liters) NaOH][conc. of standardized 0.1 M NaOH]
Once the number of moles of K+ in the weighed sample of green salt has been determined, the mass of K+
can be calculated.
mass of K+ in sample = (moles K+ )(39.10 grams/mole)
Since the mass of the green salt sample used in the experiment is known, the percentage of potassium in the
salt is given by the equation:
% K = [(mass of K+ in sample)/(mass of sample of green crystals)] x 100
-9-
Determining the Percentage of Iron in KxFey(C2O4)z . nH2O
Using Cation Exchange
An examination of the following equation
x R-SO3- H3O+ + x K+ + Fey (C2O4)z(aq)-x -----> R-SO3- K+ + x H3O+ + Fey (C2O4)z(aq)-x
indicates that the solution which elutes from the column contains the acid, x H3O+, and Fe(C2O4)z-x. When
this acid is titrated with standardized NaOH, the reaction that occurs first is given by the following equation:
H3O+ + OH- ---------> 2 H2O
After all of the acid is neutralized in the titration, addition of more NaOH results in the reaction represented as:
Fe(C2O4)z-x + 3 OH- -----------> Fe(OH)3 (ppt) + z C2O4-2
The ferric hydroxide precipitates from the solution as a reddish-brown, gelatinous precipitate. The above
equation indicates that three moles of hydroxide ion are required to react with each mole of iron ion in the
salt. Thus it follows that:
-/3 = moles of Fe in the sample
moles of OHmoles of Fe = [Volume (in liters) NaOH][conc. of standardized 0.1 M NaOH] / 3
The mass of iron (Fe) is given by the equation:
mass of Fe = (moles of Fe)(55.85 grams/mole)
Therefore the percentage of iron (Fe) in the sample is found by:
% Fe = [(mass of Fe in sample)/(mass of sample of green crystals)] x 100
The Titration Curve
When the elutant from the ion exchange column is titrated with 0.1000 M NaOH using a pH meter to follow
the course of the reaction, a titration curve looking something like the curve below is obtained:
pH
V1
V2
ml of NaOH used
-10Two titration endpoints are obtained: the first, after the addition of V1 ml of NaOH; & the second, after V2 ml
of NaOH have been added. The first endpoint represents the completion of the neutralization of the hydronium
ion (H3O+), & the second endpoint represents the completion of the precipitation of ferric hydroxide Fe(OH) . Thus V (in units of liters) represents the OH- necessary to neutralize the hydronium ion (H O+)
3
1
3
eluted and V2 - V1 (converted to liters) represents the OH- necessary to completely precipitate the Fe(OH)3.
Thus from a single pH titration curve of a weighed sample of K Fe (C O ) . nH O that has been passed
x
y
2 4 z
2
through a cation exchange resin, it is possible to determine both the % K and the % Fe in the compound.
Objectives:
1. To learn the principles and practice of using an ion exchange column.
2. To determine the % K and % Fe in KxFey(C2O4)z . nH2O using an ion exchange column.
3. To learn the principles and practice of using a pH meter.
Apparatus:
Ion exchange column (Bio Rad Econo-Column, cat. # 737-0716) packed with about 1.1 grams of Bio-Rad AG
50 W-X2 100-200 mesh analytical grade cation exchange resin (cat. # 142-1241), 10-ml graduated cylinder,
pH Hydrion paper, electronic top-loader balance, analytical balance, 50-ml beaker, buret, buret funnel, buret
clamp & support stand, magnetic stirrer, magnetic stir bar.
Chemicals:
Standardized (3 sig. fig.) 0.1 M, buffer solutions (pH = 4 and pH = 7) for pH meter standardization,
1.0 M HCl(aq), student prepared green crystals of the complex salt
Safety, Environmental, and Economic Concerns:
Get directions from your instructor on the proper disposal of waste solutions.
Notes on Experimental Procedure:
1. Mount the ion exchange column vertically on a ring stand using a utility clamp. Column should always
be filled with liquid up to or above the top of the resin so air pockets cannot form in the resin bed.
2. When you receive the column it will contain an acid solution which has been added to insure that all of the
resin is in its acid form. It is absolutely necessary to rinse out all of this acid or incorrect results will be
obtained. Never allow the liquid level to fall below the top of the resin, though!
3. During the time the solution of green salt and rinses are moving down the column, the K+ ions from the
green salt are exchanging with the hydronium ion (H3O+) of the resin.
4. As the HCl solution moves down the column, the H3O+ ions of the acid exchange with the K+ ions bound
to the resin (regenerating the resin to the acid form so that the column will be ready for the next student).
5. Be alert!!! The first equivalence point should come before 10 ml of NaOH have been added, and the
second equivalence point should come before 20 ml of NaOH have been added.
-11Experimental Procedures:
Note: This experiment should be started at the very beginning of the laboratory period in order to complete
it in time.

Obtain an ion exchange column and mount it on a ring stand. It is important to make sure that the
resin bed is filled with liquid at all times.

Using a clean, dry 10-ml graduated cylinder, rinse the column by pouring 4 ml of distilled water into
the column and collect the liquid that elutes from the column in a small beaker.

Using a piece of wide-range pH paper, test the pH of the solution that first elutes from the column to
make sure that it is distinctly acid (pH << 7) If it is not acidic, immediately inform the teacher.

Allow the level of rinse liquid in the column to fall just to the top of the resin.

Repeat the above rinse procedure two more times.

When the level of water in the third rinse has dropped to the level of the top of the resin, test a drop
of eluate with pH paper to confirm that the pH is about that of distilled water. If the eluate is still
acidic, continue to rinse until the eluate is about the pH of distilled water.

There will be some waiting time accompanying the rinse procedure. Make good use of this time by
looking ahead and preparing everything that will be needed for the day's work.

Using the electronic top-loader balance, weigh about 0.16 grams of the green crystal salt in a 50-ml
beaker. Use unheated green crystals - not the anhydrous sample from the first analysis experiment.
Make sure that the sample mass does not exceed 0.165 grams! Take the weighed sample and another
clean, dry 50-ml beaker to the analytical balance. Weigh the empty beaker to the nearest 0.0001
gram. Remove the beaker and pour the preweighed sample into the weighed beaker. Weigh the
beaker plus sample. Record these masses. [Remember that no chemical is ever added to or removed
from a container while the container is in the balance compartment. This is to avoid damaging the
balance pan as well as to prevent chemical spills in the weighing stage.]

Using a 10-ml graduated cylinder, add 4 ml of distilled water to the sample of green salt and gently
swirl the beaker until the salt is completely dissolved.

Place a clean, dry 150-ml beaker under the ion exchange column and quantitatively transfer the
solution of green salt to the column. Collect the elutant in the 150-ml beaker.

Rinse the 50-ml beaker that held the green salt solution with a 4-ml aliquot of water and when the
level of liquid in the column has dropped to the top of the resin, pour this rinse water into the
column. [The 150-ml beaker should remain under the column during this entire rinse process in
order to collect the elutant.]


Repeat this rinse procedure with two more 4-ml aliquots of distilled water, waiting each time until
the liquid level in the column has dropped to the top of the resin before adding more liquid to the
column.
Set up a buret that has been thoroughly cleaned. Set up the pH meter assembly, including magnetic
stirrer, stir bar, and distilled water bottle.
-12-

Standardize the pH meter with buffers of pH = 4 and pH = 7.

Obtain about 60 ml of the standardized 0.1 M NaOH solution in a clean, dry 100-ml beaker, and after
rinsing the buret with a 5-ml aliquot of the NaOH solution, fill the buret with the base solution.

Prepare for titration by expelling air bubbles from both the barrel and the tip of the buret.

After the last rinse has eluted from the ion exchange column, transfer the beaker containing the
elutant to the surface of the magnetic stirrer to be used in the pH meter titration.

[Before beginning the titration, pour 10 ml of 1.0 M HCl solution into the ion exchange column and
place a beaker under the column to collect the eluate. This will regenerate the column for the next
student who has to use this column.]

[Note: do not mix this elutant with the elutant to be used in the titration.]

Adjust the level of liquid in the beaker by adding distilled water so that the bulb of the pH electrode
is completely immersed in the solution.

Proceed to titrate the solution with the standardized 0.1 M NaOH. Start with 1-ml increments, then
reduce the volume as the equivalence points are reached. If the initial titration is performed
correctly, a second titration will not be necessary. Prudence is advised! Keep the volume increments
small until after the second endpoint has been passed.

Continue to titrate to about 15 ml past the second equivalence point (pH should not change when 2
ml of base is added).

Repeat the procedure for 2 more samples.

In the lab report, be sure to include a graph that plots pH data versus volume (in ml) of the NaOH
solution. Be sure that the x-axis is sufficiently expanded to allow estimations to the nearest 0.01 ml.
Draw the best smooth curve through the points. [Nothing is gained in this experiment by expansion
of the y-axis beyond the width of the graph paper.

For calculations, be sure to determine both endpoints (V1 and V2).

Calculate the average percentage iron for the 3 trials.
-13-
Determining the Percentage of Iron in KxFey(C2O4)z . nH2O
Using Cation Exchange
Data Table:
Trial 1
Trial 2
Trial 3
mass of beaker and green salt sample (in grams).....__________
__________
__________
mass of beaker (in grams)........................................ __________
__________
__________
mass of sample (in grams)....................................... __________
__________
__________
initial pH of green salt solution ............................... __________
__________
__________
V1, volume required for 1st equivalence point ......... __________ __________
__________
mass of potassium in sample (in grams).................. __________
__________
__________
percentage of K in sample ....................................... __________
__________
__________
V2, volume required for 2nd equivalence point ........ __________ __________
__________
V2 - V1, volume of NaOH that reacted with iron ...... __________ __________
__________
mass of iron in sample (in grams)............................. __________
__________
__________
percentage of iron in sample ..................................... __________ __________
__________
average percentage of iron in sample ................................................ __________
-14-
Determination of the % Oxalate in an Iron Oxalato Complex Salt
Introduction:
The percentage oxalate in KxFey(C2O4)z . nH2O will be determined by titrating a solution containing a known
mass of green salt with a standardized (3 sig. fig.) solution of 0.01 M KMnO4. The mass and percentage of
oxalate in the sample can be determined by measuring the volume, V, of KMnO4 of molarity, M, required to
completely oxidize the oxalate ion (C2O4-2). KMnO4 oxidizes C2O4-2 in acid solution to release CO2 :
16 H+ + 5 C2O4-2 + 2 MnO4- ------------> 10 CO2 + 2 Mn+2 + 8 H2O
The stoichiometry of this reaction requires an oxalate/permanganate ion ratio of 5 : 2, or stated differently, 5
moles of oxalate ion are chemically equivalent to 2 moles of the permanganate ion.
In this reaction it follows that the number of moles of oxalate ion in the sample is:
moles C2O4-2 = (V x M) x 5/2
V = volume (in liters) of KMnO4 required to react with C2O4-2; M = molarity of KMnO4 solution
The mass and percentage of oxalate in a weighed sample of the green salt is calculated from:
mass of oxalate in sample = (5/2)(V)(M)(Formula Weight of oxalate)
[Formula Weight = molar mass of the oxalate ion = 88.02 grams/mole]
% C2O4-2 = [(mass of C2O4-2)/(mass of green salt sample)] x 100
With the completion of this experiment, the percent composition of the green salt, KxFey(C2O4)z . nH2O, has
been completely determined experimentally. The simplest (empirical) formula (x, y, z) can be calculated
from the % composition. Once the formula is known it is possible to calculate the percent yield obtained in
the original synthesis reaction.
Objectives: To determine the percentage of oxalate ion in KxFey(C2O4)z . nH2O.
Apparatus:
Analytical & top-loader balance, spatula, 250-ml Erlenmeyer flask, buret funnel, buret, buret
clamp, ring support stand, 25-ml graduated cylinder, 50-ml beaker.
Chemicals:
student-prepared sample of KxFey(C2O4)z . nH2O, 0.0100 M KMnO4 solution, 85% H3PO4
solution (concentrated), 6 M H2SO4
Safety, Environmental, and Economic Concerns:
All waste solutions from this experiment may be poured down the drain and flushed with water.
Notes on Experimental Procedures:

The oxidation state of iron in KxFey(C2O4)z . nH2O is +3, normally the highest value for iron. Thus
the KMnO4 does not oxidize the iron in this experiment. However, the presence of the ferric ion
imparts a yellow color to the solution. The H3PO4 forms a colorless phosphate complex with the iron,
making it easier to detect the color change which occurs at the equivalence point.

The reaction of the permanganate ion with the oxalate ion is rather slow at room temperature. The
purpose of
-14-
Experimental Procedure:
1. Weigh a sample, to the nearest 0.0001 gram, of about 0.125 grams of the green salt crystal into a 50 ml
beaker.
2. Transfer the green crystals to a 250 ml Erlenmeyer flask, using four 15-ml aliquots of distilled water,
followed by a 6 ml aliquot of 6 M H2SO4. Add 1 ml of 85% (concentrated) H3PO4 to the sample.
3. Heat the solution to just below the boiling point (90-95º C).
4. While the first solution is being heated, thoroughly rinse a clean, dry buret with three small portions of the
standardized potassium permanganate (KMnO4) solution.
5. Fill the buret with the standardized KMnO4 solution, expel air bubbles from the buret tip and take an
initial buret reading.
6. Remove the Erlenmeyer flask from the heat source and titrate the green salt solution with the KMnO4
solution. Titrate quickly so that the reaction occurs at a nearly constant temperature.
7. Repeat this procedure with a second and third sample (be sure to heat this solution as well).
8. The equivalence point is reached once the solution has turned a very faint pink color.
Data Table:
Trial 1 Trial 2
Trial 3
mass of sample (in grams)...................................
_________
_________
_________
volume of KMnO4 required (in ml).........................
_________
_________
_________
moles of KMnO4 required ..................................
_________
_________
_________
moles of C2O4-2 in the sample .............................
_________
_________
_________
mass of C2O4-2 in the sample ..............................
_________
_________
_________
percentage of C2O4-2 in the sample .......................
_________
_________
_________
average percentage of C2O4-2 in the green salt complex ............
______________
Note on preparing the FeCl3 solution:
Into a 600 or 800 mL beaker, add 333.3 g of FeCl3 6 H2O. Slowly add about 7 mL of 6M HCl
(HCl is added to prevent the formation of hydrolysis products; the FeCl3 6 H2O is actually
dissolved by the distilled water). Caution should be exercised during the solution preparation since
the process is exothermic. Slowly add distilled water to bring the volume to approximately 400 mL.
Use magnetic stirrer to mix contents. If a reddish precipitate forms, add a minimum amount of 6M
HCl to clear the solution (do not add excessive amounts of HCl since the solution will become too
acidic resulting in poor product yields). Quantitatively transfer the contents of the beaker to a 500 mL
volumetric flask and bring volume to 500 mL mark with distilled water and/or 6M HCl (a total of 3540 mL of HCl may be necessary to produce a clear solution). The resulting solution will be deep red
or reddish orange in color.